Contrary to other recent reports, Pauling's original electronegativity equation, applied as Pauling specified, describes quite accurately homolytic bond dissociation enthalpies of common covalent bonds, including highly polar ones, with an average deviation of +/-1.5 kcal mol(-1) from literature values for 117 such bonds. Dissociation enthalpies are presented for more than 250 bonds, including 79 for which experimental values are not available. Some previous evaluations of accuracy gave misleadingly poor results by applying the equation to cases for which it was not derived and for which it should not reproduce experimental values. Properly interpreted, the results of the equation provide new and quantitative insights into many facets of chemistry such as radical stabilities, factors influencing reactivity in electrophilic aromatic substitutions, the magnitude of steric effects, conjugative stabilization in unsaturated systems, rotational barriers, molecular and electronic structure, and aspects of autoxidation. A new corollary of the original equation expands its applicability and provides a rationale for previously observed empirical correlations. The equation raises doubts about a new bonding theory. Hydrogen is unique in that its electronegativity is not constant.

[reaction: see text] In contrast to 1,3-butadiene, the textbook example of "conjugation stabilization", G3(MP2) calculations yielding the enthalpy of hydrogenation Delta(hyd)H(298) of 1,3-butadiyne indicate that it is not stabilized by the conjugated configuration of its triple bonds. Differences between ethylenic and acetylenic pi bonds are examined in the light of CAS-MCSCF calculations on 1,3-butadiene and 1,3-butadiyne.

By analogy to conjugated polyenes, conjugative stabilization of polyynes with the -CC-CC- group might be expected to be substantial. On the contrary, consistent with our recent report of a surprising lack of conjugative stabilization in butadiyne, we find by G3(MP2) calculations and by comparisons with available experimental data from these and other laboratories that the ground-state stabilization of conjugated polyynes is in fact quite small, amounting to <1 kcal mol(-)(1). By similar calculations, the 2,4-pentadiyn-1-yl radical shows no enhanced stabilization relative to 2-propyn-1-yl radical, despite the potential stabilization of the odd electron by two conjugated triple bonds and unlike the behavior of 2,4-pentadien-1-yl radical. The thermochemistry of straight-chain alkynes and polyynes is very self-consistent. Enthalpies of hydrogenation, leading to enthalpies of formation, are predictable with a high degree of accuracy (absolute mean deviation = +/-0.39 kcal mol(-)(1) vs theoretical values and +/-0.52 vs experimental) from three molecular structure enthalpies and one conjugation stabilization parameter.

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In this work, the aromaticity of pyracylene (2) was investigated from an energetic point of view. The standard enthalpy of hydrogenation of acenaphthylene (1) to acenaphthene (3) at 298.15 K was determined to be minus sign(114.5 +/- 4.2) kJ x mol(-1) in toluene solution and minus sign(107.9 +/- 4.2) kJ x mol(-1) in the gas phase, by combining results of combustion and reaction-solution calorimetry. A direct calorimetric measurement of the standard enthalpy of hydrogenation of pyracylene (2) to pyracene (4) in toluene at 298.15 K gave -(249.9 plus minus 4.6) kJ x mol(-1). The corresponding enthalpy of hydrogenation in the gas phase, computed from the Delta(f)H(o)m(cr) and DeltaH(o)m(sub) values obtained in this work for 2 and 4, was -(236.0 +/- 7.0) kJ x mol(-1). Molecular mechanics calculations (MM3) led to Delta(hyd)H(o)m(1,g) = -110.9 kJ x mol(-1) and Delta(hyd)H(o)m(2,g) = -249.3 kJ x mol(-1) at 298.15 K. Density functional theory calculations [B3LYP/6-311+G(3d,2p)//B3LYP/6-31G(d)] provided Delta(hyd)H(o)m(2,g) = -(244.6 +/- 8.9) kJ x mol(-1) at 298.15 K. The results are put in perspective with discussions concerning the "aromaticity" of pyracylene. It is concluded that, on energetic grounds, pyracylene is a borderline case in terms of aromaticity/antiaromaticity character.

Determination of enthalpies of formation, now well into its second century, continues to be an active research field. Classical combustion thermochemistry, known by Lavoisier, is carried out with precision in several laboratories, though usually on the microscale, as appropriate to the small quantities of rare or unstable species preparative chemists are able to win and purify. Nonclassical methods such as differential scanning calorimetry and proton emission techniques are practiced. Enthalpy estimation based on additivity has been brought to an improved level of accuracy, and its basis in molecular structure has been examined with the goal of achieving maximum simplicity. Discrepancies between experimental results and additive estimates due to 'special effects' have brought about a considerable amount of causative speculation in the literature. Quantum mechanical methods have enjoyed increased proliferation through new methods of finding enthalpies of formation and other thermochemical and molecular properties such as heat capacity and entropy. Powerful basis set and configuration interaction software is available within the Gaussian c suites of programs. New levels of accuracy, in the kilojoules per mole range, have been achieved by Wn methods, and wider generality is enjoyed by methods based on density functional theory. New tabulation methods have been introduced that use computer error estimation procedures to root out flawed experimental results and increase overall reliability of the data one selects from the compilation. C 2012 John Wiley & Sons, Ltd.in the gaseous state at 298 K, arrive at an enthalpy level that we shall denote H 298 C(g) 6+ + 4H + (g) + 10e − = CH 4 (g)Perhaps more familiar to chemists is the enthalpy of a molecule in the standard state referred to elements also in their standard states. For methane, C(s, gr) + 2 H 2 = CH 4 (g)where (s, gr) indicates the graphitic form of solid carbon. The enthalpy difference between the molecule and its elements, all in the standard states, 6,7 is its enthalpy (frequently called 'heat') of formation f H • . We define the enthalpy of formation of any element in its standard state as zero at all temperatures. This definition works because elements are not converted from one to another in chemical reactions.

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